Putrid-Repeat t1_iu5d8ql wrote on October 28, 2022 at 5:48 PM

Reply to comment by SmorgasConfigurator in Is an ionic bond really stronger than a covalent bond??? by jeez-gyoza

One thing I would point out is that the difference in energies between the start and end states are the same no matter the situation. The the examples you give are where the bonds can more easily be broken because the intermediate states are more easily reached. I.e. before the O2 is separated out most first have the bonds Bricker and both O2 would be so reactive that in the time they are separate they would just bond back to one another. This requires a lot of energy i.e. to make something catch fire you need to first provide enough energy to do this, usually with something very hot. In our bodies the intermediate state where they are not bound or loosely bound is stabilized by specific molecules in our body making it easier to break the bond in O2. Since this intermediate state is stabilized you don't need the large initial energy to initiate the reaction (some energy is still needed). However, the energy produced is still the same as it depends on the final and initial states only. Edit: while this is true often there is more than one reaction happening such as enthalpy of solvation, solvent- solvent integrations, and solvent solute interactions.

That said the bond energy of NaCl is 787 kj/Mol and O2 is 498 kj/mol.

Is point out as well your teacher is not correct. Ionic bonds are stronger. But, in an aqueous state they can much more easily be solvated due to water stabilizing the intermediates. Ionic bonds since they do not share electrons, once separated are very easily stabilized. In covalent bonds this is generally not the case.

SmorgasConfigurator t1_iu5gtzn wrote on October 28, 2022 at 6:12 PM

I disagree. That energy you quote for NaCl assumes the two ions are pulled apart in vacuum. If we then move the respective ion into water, there is a favourable solvation free energy. So the free energy of the final state of separated NaCl in aqueous solution is different than had it been done in vacuum. Any intermediate state stabilization, that only alters the kinetics, is not relevant.

The O2 example is, I admit, a bit different in the details. But in vacuum, if I do a homolytic cleavage of the covalent bond in O2, the two radical oxygen atoms of the final state are far from stable. So any other reactant that can combine with the two oxygen atoms make that final state stabler. The tricker part is that the path from start to finish include at least one transition state. In biology and chemistry, we can alter the kinetics of this cleavage, while leaving the final state the same, by adjusting any catalytic component. That is a more subtle point. My argument is much simpler. In vacuum, O2 is a strong bond because the final state after dissociation is unstable, while in the example system, O2 separates more readily mostly because the dissociation does not generate two free radical atoms, but one where the atomic oxygen binds to iron.

Hence, the strength of a bond must either be defined with a common reference state (vacuum typically as I mention), but then bond strength is a more abstract quantity and less informative of practical questions of stability, robustness etc, or we consider the problem in full, i.e. the stability of initial and final states are part of the analysis.

Putrid-Repeat t1_iu5zxo5 wrote on October 28, 2022 at 8:22 PM

So we are on the same page for the most part here.

But for the dissolution of salt you still have the bonds or really lattice energy to overcome. The totally enthalpy is negative as you make most of the energy required to break the bond during solvation but it still requires that energy, i.e. it's endothermic. This is what I alluded to in the multiple reactions. Total enthalpy is the sum of the bond/ lattice energy + the sum of solvent-solvent attraction energy - the sum of the solute- solvent energy.

Maybe I'm missing something but you still use the standard bond enthalpy for this equation. I.e. the sum of all the reactions enthalpy. Curious on your thoughts here.

Edit: I will agree that bond strength is not especially representative to the common environments like aqueous one.